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Topic 4 Bonding . SL + HL

Topic 4 Bonding . SL + HL. Ionic bond Covalent bond Intermolecular forces Hydrogen bond Dipole-dipole attraction van der Waals’ forces Metallic bond. 4.1 Ionic bond. Ions = charged particles Ionic bond= the electrostatic bond between positively and negatively charged ions

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Topic 4 Bonding . SL + HL

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  1. Topic 4 Bonding. SL + HL Ionicbond Covalent bond Intermolecular forces Hydrogen bond Dipole-dipole attraction van der Waals’ forces Metallic bond

  2. 4.1 Ionic bond Ions = chargedparticles Ionic bond= the electrostatic bond between positively and negatively charged ions Ionic compound= a compound built from ions, i.e. a salt

  3. Atoms want to get a “Noble gas electron configuration”. One way to get that structure is to throw away the valence electron or steal some electrons to get a full outer shell.

  4. Sodium gives one electron to chlorine and both will have noble gas configuration. Which noble gas? Sodium ion: positive electric charge. Cation. Chlorideion: negative electric charge. Anion.

  5. Ioniccrystals • + charge particles and – charge particle attracts each other in three dimension and builds up a lattice/crystal. Strong electrostatic forces in three dimensions. Each cation is surrounded by anions and vice versa.

  6. Ions • Group 1: H+, Li+, Na+, K+, Rb+, Cs+, Fr+ • Group 2: Be2+, Mg2+, Ca2+, Sr2+, Ba2+ • Group 3?/13: B3+, Al3+, Ga3+ • Group 6?/16: O2-, S2-, • Group 7?/17: F-, Cl-, Br-, I-

  7. Naming compounds • Positive ions have the name of the atom: Sodium ion • Negative ions have the name of the atom (or almost) + the ending –ide: Chloride ion • Sodium + chlorine  Sodium chloride • Lithium + oxygen Lithium oxide

  8. Formula ofioniccompounds Write the chemicalformulaof the compoundsformedbetween the positive and negative ionsabove. Write the nameof the ioniccompounds

  9. Transition metals • Transition metals can often form more than one ion, e.g.: • Fe2+ and Fe3+ • Cu+ and Cu2+ • HL: colouredbecauseof d-orbitals

  10. Some common polyatomic ions • Nitrate NO3- • Hydroxide OH- • Sulphate SO42- • Carbonate CO32- • Hydrogen carbonate HCO3- (Bicarbonate) • Phosphate PO43- • Ammonium NH4+ Ca(OH)2

  11. Formula ofioniccompounds-2 Write the chemicalformulaof the compoundsformedbetweenthe positive and negative ionsabove. Writethe nameof the ioniccompounds

  12. When can we expect an Ionic bond? • The quick rule: Metal + non-metal => Ionic compounds (Salts) CuSO4 (s)

  13. Electronegativity • Electronegativity is a measure of an atoms power to attract electrons. • On the right side in the periodic table (group 7,6,5) the atoms have high values, they attract electrons readily. Best is Fluorine, e-neg 4. • On the left side the values are low. Low ability to attract electrons.

  14. Electronegativityvalues FONClBrISCH

  15. Ionic bond or not- calculate difference in electronegativity • If you want to do a more “precise” estimation you can calculate it with the help of electronegativity values. • If the difference in electronegativity > 1,7 then you can say it's an Ionic bond.

  16. Ionic bond or not?Use electronegativityvalues in the Chemistry Data booklet • NaCl • MgO • AlCl3 • SiO2 • Ca3N2

  17. Typical properties of Salts Hard, brittle, • Conduct electricity in solution or melted • High melting points => Strong bond • Hydration of Ion in Water solution

  18. 4.2 Covalent bond Electron pair bond, molecular bond • If the De-neg < 1,7 the bond is considered to be covalent • Often between non-metals • Polar or non-polar

  19. In a covalent bond the atoms share electrons with each other to get a “Noble gas electron configuration” • The bond has a direction (one atom to an other atom) • One bond consists of two electrons, an electron pair O H H

  20. Single bond: the two atoms share two electrons (1 pair) Double bond: the two atoms share four electrons (2 pairs) Triple bond: the two atoms share six electrons (3 pairs)

  21. Lewis structures • all valence electrons marked by dots or lines. Draw Lewis structures for: • F2 • NH3 • CO2 • N2 • C2H4

  22. Number of bonds bond lengths and bond strengths • As the number of shared electrons increase (single to triple) the bond lengths shortens and the bond energy increase

  23. Non-polar covalent bond • In, for example, H2 the two electrons in the bond are shared equally between the two hydrogen atoms • H-H De-neg =0 • The electron distribution is symmetrical

  24. Polar covalent bond • If two different atoms form a covalent bond there will be a difference in De-neg • The atom with highest electronegativity will have the electrons closer; they don’t share equally • Unsymmetrical electron distribution

  25. Whichmoleculescontains at leastone polar covalent bond? F2 HF NF3 SiF4 • Which bond in the pairs belowhave the highest polar character? • C-O, N-O • H-O, S-O • H-O, H-S • Se-S, Se-F

  26. Ionic, polar or non-polar covalent bond? • % ionic character of a bond: 0-90% (there are no 100% ionic compounds) H-Cl Na+Cl- Cl-Cl

  27. Dative covalent bond • In a “normal” covalent bond both atoms contribute with electrons to the bond. • Sometimes only one atom contributes with both electrons (the electron pair) to the bond • Then the covalent bond is called a dative covalent bond

  28. Examples of dative bonds • H2O + H+ H3O+ • NH3 + H+ NH4+ • C (4 ve-)+ O (6 ve-)  CO

  29. VSEPR • Valence Shell Electron Pair Repulsion • Shape and bond angles • Determine the molecules structure, the shape in 3 dimensions • Structure around a given atom is determined principally by minimising electron-pair repulsion • Bonding or non-bonding pairs will be as far apart as they can.

  30. Linear • Two negative centers • 180o • E.g. CO2

  31. Trigonal planar (flat) • Three negative centers • 120o • E.g. BF3

  32. Tetrahedral arrangement • Four negative centers • 109.5 o

  33. Method • Draw Lewis structure • Count electron pairs, minimise the repulsion • Positions of the atoms • Name of the structure

  34. Ammonia, NH3 Tetrahedral, Trigonalplanar 107o, one non-bonding electron pair • Non-bonding pair (lone pair) takes more space => reduce bond angels

  35. Water, H2OTetrahedral, Non-linear (bent) 104oTwo non-bonding electron pairs Non-bonding pair (lone pair) takes more space => reduces bond angles

  36. Non-polar and polarmolecules (dipoles) • Based on bond polarity and molecular shape • May predict how a molecule will behave with other compounds • Polar molecule = Dipole

  37. A polar molecule (a dipole) • Must have polar covalent bonds. • Look at the difference in electronegativity. FONClBrISCH AND • Unsymmetrical shape according to charge distribution. • Otherwise it will be a non-polar molecule (NOT a dipole)

  38. Dipole or not? • H2O • CO2 • NH3 • CH4 • CH3OH

  39. Allotropes • Some elements can be found in different forms • E.g: • Carbon: Diamond, graphite, C60fullerene • Oxygen, Ozone See PPT: Carbonallotropes

  40. Silicon • Metalloid, Semiconductors, non-metallic structure • Similar structure as diamond

  41. Silicon dioxide • SiO2Silica, giant structure similar to diamond, quarts • Silicates, NaSiO4, tetrahedrical, silicon-oxygen single bond

  42. 4.3 Intermolecular forces • Holds molecules together in liquids or solids (No Intermolecular forces between gaseous particles) • Weaker than covalent and ionic bonds • Hydrogen bond (Quite strong) • Dipole-dipole (Middle weak) • van der Waal’s forces (~ London dispersion forces) (Very weak) • Accounts for differences in aggregatio state and solubility

  43. van der Waal’s forces • “Vibrations” in the electron cloud => Temporary dipoles. • A temporary dipole in one molecule can induce a temporary dipole in another molecule • Exist between all molecules

  44. van der Waal’s forces, cont. The strength increases with molar mass of the molecule/atom E.g. He b.p 4 K Xeb.p. 165 K Only effective over short range so the molecule “area” is also important. E.g: Pentane, C5H12, b.p. 309 K Dimethylpropane, (CH3)4C b.p. 283 K

  45. Dipole-dipole bond • Electrostatic attraction between molecules with permanent dipoles • Stronger than van der Waals bond Hydrogen chloride M= 36,5 g/molb.p. 188 K Fluorine M= 38 g/molb.p. 85K

  46. Hydrogen bond • In molecules that contain Hydrogen bonded to Oxygen, Nitrogen or Fluorine (high electronegativity and non-bonding electron pair) • Stronger than dipole-dipole bonds • Interaction of the non-bonding electron pair in one molecule and hydrogen (with high positive charge) in another molecule.

  47. Whichintermolecular bond? • H2O b.p.= 100oCH2S b.p.= -61oC • NH3b.p.= -33oCPH3b.p.= -88oC

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